Electron configurations 136 periodic table trends

The elements Lr-112 are expected (nonrelativistically) to be d-block elements because they are expected to involve the filling of the 6d orbital. However, relativistic calculations have shown that rutherfordium prefers a 6d 7p electron configuration rather than the 6d2 configuration expected for nonrelativistic behavior and a simple extrapolation of periodic table trends. This prediction also implies that RfCLt should... 

The use of the continuation of the periodic table, the predicted electronic configurations, and the trends which become obvious from the calculations plus the semiempirical and empirical methods, allows us to offer some detailed predictions of the properties of the elements beyond lawrencium (Z = 103) (S5). Of course, these elements will first be produced at best on a one atom at a time basis, and they offer little hope of ultimate production in the macroscopic quantities that would be required to verify some of these predictions. However, many of the predicted specific macroscopic properties, as well as the more general properties predicted for the other elements, can stiU be useful in designing tracer experiments for the chemical identification of any of these elements that might be synthesized. 

Complete the concept map using the following terms electronegativity, electron configuration, periodic trends, ionic radius, atomic radius, ionization energy, and periodic table. 

The placement of a new element in the Periodic Table requires knowledge of its atomic number and electronic configuration. Even though the atomic number can be positively assigned by a-decay chains, no knowledge is obtained about the electronic configuration or chemical properties of a new element fi om these physical methods. The elements are just placed in the Periodic Table by atomic number in various groups or series based on simple extrapolation of known Periodic Table trends or firom theoretical calculations and predictions of the electronic structures. It remains to the experimental chemist to attempt to validate or contradict these predictions. 

Ion formation is only one pattern of chemical behavior. Many other chemical trends can be traced ultimately to valence electron configurations, but we need the description of chemical bonding that appears in Chapters 9 and 10 to explain such periodic properties. Nevertheless, we can relate important patterns in chemical behavior to the ability of some elements to form ions. One example is the subdivision of the periodic table into metals, nonmetals, and metalloids, first introduced in Chapter 1. 

The outermost electrons, often called the valence electrons, are primarily responsible for the chemical properties of the elements. It follows that the elements in a specific group will show similar characteristic oxidation numbers (charges, also called valences) and display a trend in characteristics. Even though electron configurations were not known when the earliest periodic tables were formulated, the elements were placed by similarity of characteristics. 

The periodic table is a tremendous source of information for students who learn to use it well. In Chapter 4, we will learn to use the periodic table to predict the electronic configuration of each of the elements, and in Chapter 5, we will use it to predict outermost electron shell occupancy. The table s numeric data are used in later chapters on formula calculations and stoichiometry, and its information on chemical trends is applied in the chapters on bonding and molecular structure. 

You will relate the group and period trends seen in the periodic table to the electron configuration of atoms. 

Table 8.1 lists some key physical properties of the elements of the first transition series, taken mostly from Appendix F. The general trends in all of these properties can be understood by recalling that nuclear charge also increases across a period as electrons are being added to the same subshell, in this case, the d shell. The first and second ionization energies tend to increase across the period, but not smoothly. The energies of the 4s and 3d orbitals are so close to one another that the electron configurations of the neutral atoms and their ions are not easily predicted from the simplest model of atomic structure. 

The periodic table has been described as the chemist s best friend. Chemical reactions involve loss, gain, or sharing of electrons. In this chapter, we have seen that the fundamental basis of the periodic table is that it reflects similarities and trends in electron configurations. It is easy to use the periodic table to determine many important aspects of electron configurations of atoms. Practice until you can use the periodic table with confidence to answer many questions about electron configurations. As we continue our study, we will learn many other useful ways to interpret the periodic table. We should always keep in mind that the many trends in chemical and physical properties that we correlate with the periodic table are ultimately based on the trends in electron configurations. 

It is not necessary to learn the numerical values of the various physical properties, but the trends that they follow from left to right and from the top to the bottom of the Periodic Table are important. Thus, it is useful information that with the exception of H and He, fluorine, F (64 pm), is the smallest element and potassium, K (227 pm) is the largest element in the first 18 elements of the Periodic Table. These three physical properties are clearly determined by the electron configurations of the elements and their positions in the Periodic Table. It is just this combination of these three physical properties that is responsible for the chemical properties of the elements. 

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